Intermolecular forces (IMFs) are weak electrostatic attractions between molecules. They are not chemical bonds – they are broken during changes of state, not during chemical reactions.
| IMF Type | Cause | Strength | Present In |
|---|---|---|---|
| London Dispersion Forces (LDF) | Temporary, instantaneous dipoles from random electron fluctuations → induced dipoles in neighbours | Weakest (but increases with molar mass/electron count) | All molecules – polar and non-polar |
| Dipole-Dipole | Permanent dipoles attract each other (δ⁺ → δ⁻) | Moderate | Polar molecules only (e.g. HCl) |
| Hydrogen Bonding | H bonded to F, O, or N creates an extremely polar bond; the lone pair on F/O/N of another molecule attracts the δ⁺ H | Strongest IMF | H–F, H–O, H–N systems (e.g. H₂O, NH₃, HF) |
🔑 Why Water Has Anomalous Properties
Water's unusually high boiling point, high specific heat capacity, and ice's lower density than liquid water are all due to extensive hydrogen bonding between H₂O molecules. Each water molecule can form up to 4 hydrogen bonds (2 donor H atoms + 2 lone pairs on O).
⚠️ Examiner Trap – Naming IMFs
1. London forces exist in ALL substances – even polar ones. Don't say "non-polar
molecules only have London forces" and forget that polar molecules also have them.
2. Hydrogen bonding requires H bonded directly to F, O, or N – not just any
electronegative element. H–Cl does NOT show hydrogen bonding.
🔑 IB Terminology Note
The IB uses "van der Waals forces" as an umbrella term covering London (dispersion) forces, dipole-induced dipole, and dipole-dipole forces. Hydrogen bonding is not classified under van der Waals in the IB syllabus.
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